5 Complexometric Titrations

Dr. Meenu

Objectives:  To study the basics of complexometry and know the following  about self generated questions.

1. How does a complexometric titration work?
2. What are the advantages of complexometric titrations?
3. What are the different types of complexometric titration?
4. How does the indicator for complexometric titrations work?
5. What are the applications of complexometric titrations?

1 Complexometric Titrations: Geroild Schwarzenbach in 1945 for the first time introduced that aminopolycarboxylic acids can be used as titrants against the analyte (metal ions). These titrations involve combination of ions with organic ligand to form dissolved complex with analyte (a metal ion).

M (aq) + nL → MLn (aq)

2   Components of Complex:

2.1 Metal: Most metal ions can accept unshared pair of electrons from an anion or molecule to form coordinate covalent bonds. Usually metals accept 2, 4 or 6e^– pairs (coordination number)

2.2  Ligand:

a)      Monodentate ligands (2e^– donor) examples NH3, H2O

b)      Ambidentate ligands (2e^– donor) have more than one donor sites but coordinate through one donor site at one time examples SCN^–, CN^–

c)      Polydentate ligands (> 2e^– donor) ethylenediammine tetraacetic acid (EDTA)

d)     Chelate ligands forms ring around metal complex 2,2’-bipyridine (bpy) , ethylenediammine (en), ethylenediamminetetracetic acid (EDTA)

Figure 1: 1,2-Ethylenediamine as a chelating ligand for metal ion

Chelate effect: chelate effect is observed with polydentate ligands as they form thermodynamically more stable complexes. Though in case of reaction with monodentate and bidentate ligand H may remain same or slightly differ but the entropy is favourable for multi dentate ligands resulting in stable complex.

3 Advantages of Complexometric Titrations

•      Complexometric titrations are fast and accurate.

•      In complexometric titrations difficulty due to formation of stepwise complex can be minimized

For almost all the metal ions the coordination number is either 4 or 6. With monodentate ligands multi step reactions are required. So at any point of time number of complexes are present in a reaction mixture making it difficult to observe sharp equilibrium point.

To overcome these problems multidentate ligands are used. EDTA (H4Y) is the most commonly used ligand. EDTA has 6 coordination sites (4 O and 2 N) and forms 1:1 strong metal complex.

Disodium salt of EDTA Na2H2Y.2H2O is used for complexation due to insolubility of H4Y in water. Upon complexation stable M-EDTA chelate is formed. The successive pka values for EDTA are pK1 = 0.032 (COOH), pK2 = 1.5(COOH), pK3 = 2.0(COOH), pK4 =2.66(COOH), pK5 = 6.16(NH+) and pK6 = 10.24 (NH+). EDTA exists in 7 acid-base forms depending upon pH of solution as shown in Fig. 2.

Out of six coordination sites first two sites are relatively stronger acids and are of least importance in equilibrium calculations. The four dissociation constants are expressed as

The formation constant for complexes of various metal ions with EDTA are given in Table 1. The data indicates that except for group I metal ions ligand combines with almost all the metal ions. The lower stability of M-EDTA complex and high pH of solution must be maintained during titration.

Table 1:The stability constants for [M(EDTA)]⁺ⁿ complexes

From the above table, it is clear that formation constant for most EDTA complexes are quite large and the values are proportional to the size of positively charged cation. For hexacoordinated M-EDTA complex, the ring formed through wrapping of metal complex by EDTA is strained and excess strain is removed by drawing the oxygen ligands towards nitrogen. During this process one or two more extra sites may be created as is evident from octa coordinated Ca(EDTA)(H2O)2^-2.

4 Titration Curve:

Figure 2:-The complexometric titration curve

The titration curve for complexometric titrations is similar to curve obtained by titration of strong acid with weak base. The titration curve is divided into three regions

a)      Pre-equivalence point: where the concentration of uncomplexed metal ion is high. The dissociation of MYn−4 is negligible (region 1 of titration curve).

b)     Equivalence point: where concentration of uncomplexed EDTA is in equilibrium with MYn−4 (region 2 of titration curve).

c)      Post Equivalence point: where the concentration of EDTA is high (region 3 of titration curve).

5   Types of EDTA Titrations

5.1 Direct titrations: These titrations involve titration of metal ion with ligand directly. Some auxillary ligand such as ammonia, tartrate and citrate etc. should be employed to prevent metal hydroxide formation as the solutions are kept at basic pH for EDTA titrations.

5.2 Back titrations: Under certain circumstances the direct titrations are not feasible. These conditions include:

• Precipitation of metal in necessary pH range required for titration.
• Slow reaction between metal and ligand.
• Unavailability of suitable indicator.

An excess amount of EDTA solution is added to the metal solution, buffer is added to maintain the pH of the solution. The resulting solution is heated and cooled at room temperature. The excess amount of EDTA is back titrated with zinc sulphate or chloride.

5.3 Substitution titrations: The titrations where end point is not sharp or metal-indicator complex is not formed. These titrations are also suitable where metal EDTA complex is more stable than Mg-EDTA and Ca-EDTA complex.

5.4 Alkalimeteric titrations: These titrations are based on principle of redox titrations where liberated hydrogen from the reactions of heavy metal ion with disodium salt of EDTA is titated with base like NaOH without any buffer.

6  End Point Detection

•      Metallochromic indicators are used for end point detection.

•      Colored organic compounds form complexes with metal ions

•      Stability of M-EDTA complex should be higher than M-Ind complex

•      Common metallochromic indicators are as listed in Table 2

Table 2: Indicators for M-EDTA complexometric titration

• Erichrome Black T (EBT)
• EBT is useful for titrations in pH range 7-10 which includes complexation of calcium, magnesium and zinc.
• EBT is triprotic and has a different structure at varying pH values.
• Between pH 7 and 10 it exists as the dianion (HIn2-) and is blue in colour and forms red colour complexes with metal ion.

7 EDTA Titrations: Selectivity

Wide variety of metal ions form complexes with EDTA, thus for determination of particular metal in presence of other interferences, masking agents are added to bind interferences. After the initial titration, releasing agents are added to proceed in forward direction. Following methods are used for masking

7.1 Precipitation: During the titration of hard water co-precipitation of calcium and magnesium can be avoided if magnesium ions are precipitated as hydroxides at pH 12 and calcium alone can be detected in the initial reaction.

7.2 Oxidation state: Transition metal ions exist in variable oxidation state. The stability constant for one oxidation state may differ from other oxidation state. As in case of iron, Fe(III) forms stable complexes with EDTA than Fe(II).

7.3 pH: During EDTA complexation as the acidity increases, the equilibrium of reaction shifts towards left direction. So high pH is maintained throughout the reaction using a buffer solution. The calcium and magnesium ions are determined at high pH, however low pH is maintained for the determination of Fe(III).

M⁺ⁿ      +  H2Y^-2         MY^(n-4)-    +            2H⁺

7.4 Masking agents for precipitation: Masking agents are added to remove the interference of metal ion encountered during the titration of particular ion for example cyanide is added to remove the interference of Zn⁺² and Cd⁺² during the titration of Ca⁺² and Mg⁺² ions while performing the simultaneous titration with EDTA. Suitable demasking agents are used to demask the previously masked metal ion as in case of Zn⁺² and Cd⁺² chloral hydrate or a 3:1 formaldehyde : acetic acid solution is used.

Table 3: Common masking agents for metal ions in complexometric titrations with EDTA

8        Applications of Complexometric Titrations

8.1         Determination of Total and Calcium Hardness in Water

Theory: hardness of water is measured in terms of calcium and magnesium ions present in water. It may be temporary if carbonates and bicarbonates of calcium and magnesium are there in water and permanent if chlorides or sulphates of calcium and magnesium are there. Temporary hardness can be removed by boiling the water where as permanent hardness can be removed by ion exchange method. ‘Soft’ water has hardness < 60 mg CaCO3/L that may cause corrosion of pipes. ‘Hard’ water has hardness > 270 mg CaCO3/L that can lead to deposition of Ca and Mg minerals (limescale in water heaters, pipes etc.).

Analysis of the water: pipette out 25 mL of the sample solution into titration flask, add 4mL of pH 10 buffer, 2 drops of Eriochrome Black T indicator. Heat the solution to about 60 °C and titrate with EDTA solution until a purple color is obtained. Shake the solution then continue titrating, adding EDTA drop wise with constant shaking until blue color with no tinge of purple is obtained. This gives total Ca (II) and Mg(II) concentration.

Determination of Ca(II): pipette out 25 mL of sample solution and add 3 mL of 50% NaOH. The Mg(II) will precipitate as Mg(OH)2. Shake the solution vigorously. Add 2 drops of calcon indicator and titrate with EDTA drop wise until a drop of EDTA turns the solution blue, but purple color will appear within 20 seconds. Allow the solution to stand for 5 minutes with occasional stirring. Some of the calcium is precipitated as CaCO3, this allows the dissolution of CaCO3 to ensure correct end point. Continue titration unless blue color persistent for 20 seconds is obtained.

Calculation:

•      1 mL 0.01 M EDTA = 0.4008 mg Ca(II)

•      1 mL 0.01 M EDTA = 0.2731 mg Mg(II)

•      Volume of EDTA consumed with EBT for total hardness (Ca(II) and Mg(II)) = X mL

•      Volume of EDTA consumed with calcon indicator for Ca(II) only = Y mL

•      mg of Ca(II) = 0.4008 Y

•      mg of Mg(II) = 0.2731(X-Y)

9.2 Cation Exchange Capacity of Soil: Cations are positively charged particles such as Cu+2, Zn+2. The cations are held tightly on soil. Though the soil particles are made up of Si+4 but their replacement with Al+3 due to isomorphous substitution results into net negative charge. This charge is neutralised with cations present in the soil. The negative charges associated with isomorphous substitution are considered permanent, that is, the charges do not change with pH changes. However organic matter can have a 4 to 50 times higher pH dependent CEC per given weight than clay.

Effect of CEC on soil

Potassium and magnesium deficiency can be associated with low CEC value. Soil pH will decrease at intervals for low CEC soils.

Procedure: To 5.0 g of soil sample taken in 100 mL centrifuge tube add 50 mL of 1N sodium acetate (pH = 5) and stirred with glass rod. The suspension was digested in a near boiling water bath for 30 minutes with stirring. The suspension was centrifuged and decanted to remove supernatant liquid containing the soluble salts. The residue was washed with 1N sodium acetate with 30 minutes digestion, in the boiling water bath, followed by 1N CaCl2 washing. Excess salts were removed by washing with 99% acetone, until excess CaCl2 was removed. Calcium exchanged for all other cations in soil was removed by washing 1N sodium acetate and supernatant collected in 250 mL conical flask. 10 mL ammonia buffer was added to maintain pH at 10.0. An EBT indicator 2 drop was added and the solution was titrated against EDTA (0.1N) till a deep ocean colour was developed.

CEC (Meq exchange capacity per 100 g) = mL of EDTA *N* 100/mass of soil sample.

Bibliography

1. G. Marr and B.W. Rocket, ‘Practical Inorganic Chemistry’, University Science Books, 1999.
2. G. Pass and H. Sutcliffe, ‘Practical Inorganic Chemistry’, Chapman and Hall, London, 1968.
3. Vogel’s Textbook of Quantitative Chemical Analysis, Arthur Israel Vogel, Prentice Hall, 2000.
4. Gerold Schwarzenbach, Complexometric Titrations, Methuen, 1969.